Light Energy and Electron Configurations Chapter 5.1 Light Energy and Electron Configurations

QUANTUM THEORY Describes how electrons behave like both waves and particles Is the modern theory of the atom Properties of waves: property definition symbol SI units velocity distance traveled per second c m/s Amplitude peak height above the midline A m wavelength peak-to-peak distance λ Frequency number of peaks passing by per second ν s-1 (Hertz)

3 ways to tell a particle from a wave: wave behavior particle behavior waves interfere particles collide waves diffract particles effuse waves are delocalized particles are localized

Electron Configurations Chapter 5.2-5.3 Electron Configurations

The Bohr Model of the Atom Ground State the lowest energy state of an atom all electrons are in their lowest possible orbitals. when atoms are in the ground state, they do not emit light

The Bohr Model of the Atom Excited State state in which an atom or molecule picks up outside energy, causing an electron to move into a higher-energy orbital when the electron falls back to a lower-energy orbit, the electron emits an amount of energy (quantum) -- light

The Bohr Model of the Atom Electrons only move about in certain ENERGY LEVELS. The smaller the electron’s orbit, the LOWER the atom’s energy state, or ENERGY LEVEL. Bohr assigned a number, n, to each energy level and called this the PRINCIPAL QUANTUM NUMBER

The Heisenberg Uncertainty Principle It is fundamentally impossible to know precisely both the VELOCITY and LOCATION of a particle at the same time. By measuring one, the other is affected.

ATOMIC ORBITALS The quantum model of the atom shows that electrons do not follow orbits around the nucleus, but rather have areas of high probability where they may be found. These areas are called ENERGY LEVELS. We use QUANTUM NUMBERS to describe the location of an electron.

ATOMIC ORBITALS Each energy level is broken up into sub-levels that have different shapes. These sublevels are defined by their shapes as S, P, D, or F. Each sub-level contains a SPECIFIC NUMBER OF ORBITALS

WHAT DO THE ORBITALS LOOK LIKE?

ATOMIC ORBITALS Each orbital can hold 2 electrons that are spinning in OPPOSITE DIRECTIONS. Each sub-level has a different number of orbitals : Sub-Level # of Orbitals s 1 p 3 d 5 f 7

Reading the Periodic Table

WRITING ELECTRON CONFIGURATIONS The arrangement of an atom’s electrons is called the ELECTRON CONFIGURATION. Three rules define how electrons can be arranged in an atom’s orbitals:

AUFBAU PRINCIPLE Electrons must occupy the orbital with the lowest energy first

Increasing energy 6s 5s 4s 3s 2s 1s 6p 5p 5d 4p 4d 4f 3p 3d 2p

Pauli Exclusion Principle Orbitals can only have 2 electrons max No two electrons in the same atom can be in the same orbital, spinning the same way. To occupy the same orbital, two electrons must have opposite spins; that is, the electron spins must be paired.

Hund’s Rule Every orbital in a subshell is singly occupied with one electron before any one orbital is doubly occupied all electrons in single occupied orbitals have the same spin.

NOTATION: Complete electron configurations are written using: Numbers to represent ENERGY LEVEL (n) Letters to represent SUBLEVELS (l) Superscript numbers to represent INDIVIDUAL ELECTRONS The order in which the energy levels, sub-levels, and orbitals will be filled is represented on your PERIODIC TABLE

Orbitals and Electron Capacity of the First Four Principle Energy Levels Principle energy level (n) Type of sublevel Number of orbitals per type Number of orbitals per level(n2) Maximum number of electrons (2n2) 1 s 2 4 8 p 3 9 18 d 5 16 32 f 7   

Arrangement of Electrons in Atoms Electrons in atoms are arranged as SHELLS (n) SUBSHELLS (l) ORBITALS (ml)

Writing Atomic Electron Configurations Two ways of writing configs. One is called the spdf notation. 1 s value of n value of l no. of electrons spdf notation for H, atomic number = 1

Another way of writing electron configurations is called ORBITAL BOX NOTATION. This includes additional information showing each orbital and each electron including the spin!

Draw the orbital notation for chlorine Spdf notation: 1s2 2s2 2p6 3s2 3p5 Orbital Diagram: 1s 2s 2p 3s 3p

Draw the orbital notation for aluminum Spdf notation: 1s2 2s2 2p6 3s2 3p1 Orbital Diagram: 1s 2s 2p 3s 3p

Practice Problems: 1s22s22p3 Write the electron configuration for nitrogen. 1s22s22p3

Practice Problems: 1s22s22p63s23p64s2 Write the electron configuration for calcium 1s22s22p63s23p64s2

Practice Problems: 1s22s22p63s23p64s23d104p65s24d10 5p66s24f145d106p6 Write the electron configuration for Radon. 1s22s22p63s23p64s23d104p65s24d10 5p66s24f145d106p6

Practice Problems: 1s22s22p63s23p64s23d104p65s24d10 Write the electron configuration for Americium. 1s22s22p63s23p64s23d104p65s24d10 5p66s24f145d106p67s25f7

Kernel (Core) Notation I wish there was some faster way...

What is the kernel (core) of an atom? The kernel, or core, of an atom includes all of the nucleus of the atom and any electrons not on the outermost energy level. For electron configurations, you can use brackets and the elemental symbol of the previous noble gas to represent the electron configuration for the core or kernel of the atom. [Ne] = 1s22s22p6

What is the core notation for: Chlorine [Ne] 3s23p5 Uranium [Rn] 7s25f4

Orbitals fill in an order Lowest energy  higher energy. Adding e-’s changes energy of orbital. Full orbitals are best situation. half filled orbitals  lower energy, next best more stable.

Write the electron configurations for these elements: Titanium - 22 electrons 1s22s22p63s23p64s23d2 Vanadium - 23 electrons 1s22s22p63s23p64s23d3 Chromium - 24 electrons 1s22s22p63s23p64s23d4 (expected) But this is not what happens!!

Chromium is actually: 1s22s22p63s23p64s13d5 Why? This gives us two half filled orbitals (the others are all still full) Half full is slightly lower in energy. The same principal applies to copper.

Copper’s electron configuration Copper has 29 electrons so we expect: 1s22s22p63s23p64s23d9 But the actual configuration is: 1s22s22p63s23p64s13d10 This change gives one more filled orbital and one that is half filled. Remember these exceptions: d4, d9

Irregular configurations of Cr and Cu Chromium steals a 4s electron to make its 3d sublevel HALF FULL Copper steals a 4s electron to FILL its 3d sublevel

Exceptional Electron Configurations Some actual electron configurations differ from those assigned using the aufbau principle because although half-filled sublevels are not as stable as filled sublevels, they are more stable than other configurations.

Electron Configurations for IONS Isoelectronic with noble gases “iso-” = SAME “electronic”= ELECTRON CONFIGURATIONS When making positive ions… they can only lose valence electrons! When making negative ions… they can only gain valence elctrons!

What is the electron configuration for a sodium ion? Na+1 Number of Electrons = 10!! 1s22s22p6

Excited State Electron configurations Atoms that are excited have electrons that have been elevated to higher energy levels or subshells due to an input of energy. They will violate the Aufbau principle!! You can tell which element it is by counting the superscripts to see how many total electrons there are 

Valence Electrons Determine the chemical properties of an element bonding Number of electrons found in the atom’s outermost orbitals Only the s and p electrons of the highest energy level! Corresponds with the element’s group number Transition metals may have multiple valence electrons

Electron-Dot Structures Consists of an element’s symbol surrounded by dots The dots represent the valence electrons

Reading the Periodic Table